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Atomic Structure

Atomic Structure

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Atomic Structure

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  1. Atomic Structure Chapter 4

  2. Defining the Atom • Atom  smallest particle of an element that retains its identity in a chemical reaction • Early models of the atom were proposed by philosophers and scientists that could not observe individual atoms • Democritus • Dalton

  3. Democritus’s Atomic Philosophy • Matter is composed of extremely small particles that cannot be divided • He called the particles atoms, from the Greek word “atomos” meaning uncut or indivisible

  4. Dalton’s Atomic Theory • More than 2000 years after Democritus proposed his model of the atom, a schoolteacher named John Dalton formulated hypotheses to explain his observations of the atom • Dalton studied the ratios in which elements combine in chemical reactions

  5. ELEMENT 1 ELEMENT 2 ELEMENT 3 ELEMENT 4

  6. + +

  7. Sizing up the Atom • The radius of most atoms fall within the range of 5 x 10-11 m to 2 x 10-19 m • Despite their small size, individual atoms are observable with instruments such as scanning tunneling microscopes • Nanotechnology is becoming an essential application to medicine, communications, solar energy, and space exploration

  8. Checkpoint • What happens to atoms in a chemical reaction according to Dalton’s atomic theory? • Atoms can be combined, separated, or rearranged, but never changed into atoms of another element

  9. Structure of the Nuclear Atom • We now know that atoms are divisible into fundamental particles called subatomic particles • Three kinds of subatomic particles are electrons, protons, and neutrons

  10. Electrons • Discovered by JJ Thompson using a cathode-ray tube • Electrons are negatively charged subatomic particles

  11. Protons • Discovered by Eugen Goldstein who observed particles traveling in the opposite direction of negatively charged cathode rays • Protons are positively charged subatomic particles

  12. Neutrons • Discovered by James Chadwick • Neutrons are subatomic particles with no charge but with a mass nearly equal to that of the proton

  13. The Atomic Nucleus • Rutherford’s Gold-Foil Experiment • Yielded evidence of the atomic nucleus • Prior to this discovery, the plum-pudding model was believed to be the correct model of the atom • Rutherford concluded that all the positive charge and almost all the mass are concentrated in a small region he called the nucleus • Nucleus is the tiny central core of an atom and is composed of protons and neutrons • Electrons are distributed around the nucleus and occupy almost all the volume of the atom

  14. Checkpoint • How do negatively charged plates affect the path of cathode rays? • A negatively charged plate repels the cathode ray • What is the charge of a neutron? • Neutral

  15. Distinguishing among atoms • Atomic Number • Mass Number • Isotopes • Atomic mass

  16. Atomic Number • Elements are different because they contain different numbers of protons • The atomic number of an element is the number of protons in the nucleus of an atom of that element • Atoms are electrically neutral so the number of protons must be equal to the number of electrons

  17. Mass Number • Most of the mass of an atom is concentrated in its nucleus and depends on the number of protons and neutrons • Mass number is the total number of protons and neutrons in an atom • If you know the atomic number and mass number of an atom of any element, you can determine the atom’s composition • The number of neutrons in an atom is the difference between the mass number and atomic number

  18. Periodic Table • The periodic table gives us valuable information for each element # OF PROTONS + # OF NEUTRONS Cl MASS NUMBER 35 ATOMIC NUMBER 17 NUMBER OF PROTONS

  19. Representing Elements • Elements can also be represented by writing “element name-mass number” Sodium-22 or Carbon-13

  20. Complete the following chart: Remember: Protons = electrons; Mass number = protons + neutrons; Atomic number = protons

  21. Isotopes • Isotopes are atoms that have a different number of neutrons but the same number of protons • Because isotopes of an element have different numbers of neutrons, they also have different mass numbers • Isotopes are chemically alike because they have identical numbers of protons and electrons • The number of neutrons do not affect chemical behavior

  22. Isotopes of Hydrogen

  23. Checkpoint • How do you calculate mass number? • Mass number = protons + neutrons • What are three known isotopes of hydrogen? • Hydrogen, Deuterium, and Tritium

  24. Atomic Mass • An atomic mass unit is one twelfth of the mass of a carbon-12 atom • The isotope of carbon was assigned a mass of exactly 12 atomic mass units (6 protons and 6 neutrons) • The atomic mass of an element is the weighted average mass of the atoms in a neutrally occurring sample of the element • A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature

  25. Calculating Atomic Mass Atomic Mass = isotope 1 + isotope 2 + … Atomic mass x relative abundance Atomic mass x relative abundance

  26. Calculating Atomic Mass • Element X has two natural isotopes. The isotope with a mass of 10.012 amu has a relative abundance of 19.91%. The isotope with a mass of 11.000 amu has a relative abundance of 80.09%. Calculate the atomic mass of this element.

  27. Checkpoint • What is the atomic mass of an element? • Weighted average mass of the atoms in a naturally occurring sample of the element • What three values must be known in order to calculate the atomic mass of an element? • Number of stable isotopes • Mass of each isotope • Natural percent abundance of each isotope